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Uncut diamonds. The strength of the covalent bonds between

adjacent carbon atoms gives diamonds their hardness.

The Solid State 343

each bond. This leaves one outer electron in each carbon atom free to circulate

through the network, thereby accounting for graphite’s near-metallic luster and elec-

trical conductivity. Although each layer is quite strong, weak van der Waals forces

(Sec. 10.4) bond the layers together. As a result the layers can slide past each other

readily and are easily flaked apart, which is why graphite is so useful as a lubricant

and in pencils.

Under ordinary conditions graphite is more stable than diamond, so crystallizing

carbon normally produces only graphite. Because graphite is less dense than diamond

(2.25 g /cm^3 versus 3.51 g /cm^3 ), high pressures favor the formation of diamond. Natural

diamonds originated deep in the earth where pressures are enormous. To synthesize

diamonds, graphite is dissolved in molten cobalt or nickel and the mixture is

compressed at about 1600 K to about 60,000 bar. The resulting diamonds are less than

1 mm across and are widely used industrially for cutting and grinding tools.

0.154 nm

Figure 10.8The tetrahedral structure of diamond.

The coordination number is 4.

Figure 10.9Graphite consists of layers of car-

bon atoms in hexagonal arrays, with each

atom bonded to three others. The layers are

held together by weak van der Waals forces.

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